Acid-base reaction
An acid–base reaction is a chemical reaction that occurs
between an acid and a base. Several concepts that provide
alternative definitions for the reaction mechanisms involved
and their application in solving related problems exist. Despite
several differences in definitions, their importance becomes
apparent as different methods of analysis when applied to
acid–base reactions for gaseous or liquid species, or when acid
or base character may be somewhat less apparent. The first of
these scientific concepts of acids and bases was provided by
the French chemist Antoine Lavoisier, circa 1776
Historic acid–base theories:
Lavoisier's oxygen theory of acids:
The first scientific concept of acids and bases was provided
by Lavoisier circa 1776. Since Lavoisier's knowledge of strong
acids was mainly restricted to oxoacids, such as HNO3 (nitric
acid) and H2SO4 (sulphuric acid), which tend to contain central
atoms in high oxidation states surrounded by oxygen, and since
he was not aware of the true composition of the hydrohalic
acids (HF, HCl, HBr, and HI), he defined acids in terms of their
containing oxygen, which in fact he named from Greek words
meaning "acid-former" (from the Greek οξυς (oxys) meaning
"acid" or "sharp" and γεινομαι (geinomai) meaning
"engender"). The Lavoisier definition was held as absolute truth
for over 30 years, until the 1810 article and subsequent
lectures by Sir Humphry Davy in which he proved the lack of
oxygen in H2S, H2Te, and the hydrohalic acids. However, Davy
failed to develop a new theory, concluding that "acidity does
not depend upon any particular elementary substance, but
upon peculiar arrangement of various substances" .One
notable modification of oxygen theory was provided
by Berzelius, who stated that acids are oxides of nonmetals
while bases are oxides of metals.
Liebig's hydrogen theory of acids:
This definition was proposed by Justus von Liebig circa
1838, based on his extensive works on the chemical
composition of organic acids. This finished the doctrinal shift
from oxygen-based acids to hydrogen-based acids, started by
Davy. According to Liebig, an acid is a hydrogen-containing
substance in which the hydrogen could be replaced by a
metal. Liebig's definition, while completely empirical, remained
in use for almost 50 years until the adoption of the Arrhenius
definition.
Common acid–base theories:
The Arrhenius definition of acid–base reactions is a
development of the hydrogen theory of acids, devised
by Svante Arrhenius, which was used to provide a modern
definition of acids and bases that followed from his work
with Friedrich Wilhelm Ostwald in establishing the presence of
ions inaqueous solution in 1884, and led to Arrhenius receiving
the Nobel Prize in Chemistry in 1903 for "recognition of the
extraordinary services... rendered to the advancement of
chemistry by his electrolytic theory of dissociation".
As defined by Arrhenius, acid–base reactions are characterized
by Arrhenius acids, whichdissociate in aqueous solution to form
hydrogen ions (H+), and Arrhenius bases, which form hydroxide
(OH−) ions. More recent IUPAC recommendations now suggest
the newer term "hydronium" be used in favor of the older
accepted term "oxonium" to illustrate reaction mechanisms
such as those defined in the Brønsted–Lowry and solvent
system definitions more clearly, with the Arrhenius definition
serving as a simple general outline of acid–base character. The
Arrhenius definition can be summarised as "Arrhenius acids
form hydrogen ions in aqueous solution with Arrhenius bases
forming hydroxide ions."
The universal aqueous acid–base definition of the Arrhenius
concept is described as the formation of water from hydrogen
and hydroxide ions, or hydrogen ions and hydroxide ions from
the dissociation of an acid and base in aqueous solution:
H+ (aq) + OH− (aq)
H2 O
(In modern times, the use of H+ is regarded as a shorthand
for H3O+, since it is now known that the bare proton H+ does
not exist as a free species in solution.)
This leads to the definition that in Arrhenius acid–base
reactions, a salt and water are formed from the reaction
between an acid and a base. In other words, this is
a neutralization reaction.
acid + base → salt + water
The positive ion from a base forms a salt with the negative ion
from an acid. For example, two moles of sodium ion (Na+) from
the base sodium hydroxide (NaOH) combine with one mole of
sulfate ion (SO42-) from sulfuric acid (H2SO4) to form one mole
of sodium sulfate(Na2SO4) . Two moles of water are also
formed.
2 NaOH + H2SO4 → Na2SO4 + 2 H2O
The Arrhenius definitions of acidity and alkalinity are restricted
to aqueous solutions, and refer to the concentration of the
solvent ions. Under this definition, pure H2SO4 or HCl dissolved
in toluene are not acidic, and molten KOH and solutions of
sodium amide in liquid ammonia are not alkaline.
Solvent system definition:
One of the limitations of Arrhenius definition was its reliance
on water solutions. Edward C. Franklin studied the acid–base
reactions in liquid ammonia in 1905 and pointed out the
similarities to water-based Arrhenius theory, and Albert F. O.
Germann, working with liquid COCl2, generalized Arrhenius
definition to cover aprotic solvents and formulated the solvent
system theory in 1925.
Germann pointed out that in many solvents there is a certain
concentration of a positive species, solvonium (earlier lyonium)
cations and negative species, solvate (earlier lyate) anions, in
equilibrium with the neutral solvent molecules. For example,
water and ammonia undergo such dissociation
into hydronium and hydroxide, and ammonium and amide,
respectively:
2 H2 O
2 NH3
H3O+ + OH−
NH4+ + NH2-
Some aprotic systems also undergo such dissociation, such
as dinitrogen tetroxide into nitrosonium and nitrate, antimony
trichloride into dichloroantimonium and
tetrachloroantimonate, and phosgene into chlorocarboxonium
and chloride.
NO+ + NO3+
N2O4
2 SbCl3
COCl2
SbCl2+ + SbCl4COCl+ + Cl-
A solute causing an increase in the concentration of the
solvonium ions and a decrease in the solvate ions is defined as
an acid and one causing the reverse is defined as a base. Thus,
in liquid ammonia, KNH2 (supplying NH2-) is a strong base,
and NH4NO3 (supplying NH4+) is a strong acid. In liquid sulfur
dioxide (SO2), thionyl compounds (supplying SO2+) behave as
acids, and sulfites (supplying SO32-) behave as bases.
The non-aqueous acid–base reactions in liquid ammonia are
similar to the reactions in water:
2 NaNH2 (base) + Zn(NH2)2 (amphiphilic amide) → Na2[Zn(NH2)4]
2 NH4I (acid) + Zn(NH2)2 (amphiphilic amide) → [Zn(NH3)4)]I2
Nitric acid can be a base in liquid sulfuric acid:
HNO3 (base) + 2 H2SO4 → NO2+ + H3O+ + 2 HSO4The unique strength of this definition shows in describing the
reactions in aprotic solvents, for example in liquid N2O4:
AgNO3 (base) + NOCl (acid) → N2O4 (solvent) + AgCl (salt)
Since solvent-system definition depends on the solvent as well
as on the compound itself, the same compound can change its
role depending on the choice of the solvent. Thus, HClO4 is a
strong acid in water, a weak acid in acetic acid, and a weak base
in fluorosulfonic acid. This was seen as both a strength and a
weakness, since some substances, such as SO3 and NH3, were
felt to be acidic or basic on their own right. On the other hand,
solvent system theory was criticized as too general to be useful;
it was felt that there is something intrinsically acidic about
hydrogen compounds, not shared by non-hydrogenic
solvonium salts.
Brønsted–Lowry definition:
The Brønsted–Lowry definition, formulated in 1923,
independently by Johannes Nicolaus Brønsted in Denmark
and Martin Lowry in England, is based upon the idea
of protonation of bases through the de-protonation of acids –
that is, the ability of acids to "donate" hydrogen ions (H+)
or protons to bases, which "accept" them. Unlike the previous
definitions, the Brønsted–Lowry definition does not refer to the
formation of salt and solvent, but instead to the formation
of conjugate acids and conjugate bases, produced by the
transfer of a proton from the acid to the base. In this approach,
acids and bases are fundamentally different in behavior from
salts, which are seen as electrolytes, subject to the theories
of Debye, Onsager, and others. An acid and a base react not to
produce a salt and a solvent, but to form a new acid and a new
base. The concept of neutralization is thus absent.
According to Brønsted–Lowry definition, an acid is a compound
that can donate a proton, and a base is a compound that can
receive a proton. An acid–base reaction is, thus, the removal of
a hydrogen ion from the acid and its addition to the base. This
does not refer to the removal of a proton from the nucleus of
an atom, which would require levels of energy not attainable
through the simple dissociation of acids, but to removal of a
hydrogen ion (H+).
The removal of a proton (hydrogen ion) from an acid produces
its conjugate base, which is the acid with a hydrogen ion
removed, and the reception of a proton by a base produces
its conjugate acid, which is the base with a hydrogen ion added.
For example, the removal of H+ from hydrochloric acid (HCl)
produces the chloride ion (Cl−), the conjugate base of the acid:
HCl → H+ + Cl−
The addition of H+ to the hydroxide ion (OH−), a base, produces
water (H2O), its conjugate acid:
H+ + OH− → H2O
Although Brønsted–Lowry acid–base behavior is formally
independent of any solvent, it encompasses Arrhenius and
solvent system definitions in an unenforced way. For example,
protonation of ammonia, a base, gives ammonium ion, its
conjugate acid:
H+ + NH3 → NH4+
The reaction of ammonia, a base, with acetic acid in absence of
water can be described to give ammonium cation, an acid, and
acetate anion, a base:
CH3COOH + NH3 → NH4+ + CH3COO−
This definition also explains the dissociation of water into low
concentrations of hydronium and hydroxide ions:
H2 O + H2 O
H3O+ + OH−
Water, being amphoteric, can act as both an acid and a base;
here, one molecule of water acts as an acid, donating a H+ ion
and forming the conjugate base, OH−, and a second molecule of
water acts as a base, accepting the H+ ion and forming the
conjugate acid, H3O+.
Acid dissociation and acid hydrolysis are seen to be entirely
similar phenomena:
HCl (acid) + H2O (base) H3O+ (acid) + Cl− (base)
NH4+ (acid) + H2O (base) H3O+ (acid) + NH3 (base)
as are basic dissociation and basic hydrolysis:
NH3 (base) + H2O (acid) NH4+ (acid) + OH− (base)
CH3COO− (base) + H2O (acid) CH3COOH (acid) + OH− (base)
Thus, the general formula for acid–base reactions according to
the Brønsted–Lowry definition is:
AH + B → BH+ + A−
where AH represents the acid, B represents the
base, BH+ represents the conjugate acid of B, and A− represents
the conjugate base of AH.
Although Brønsted–Lowry calls hydrogen-containing
substances like HCl acids, KOH and KNH2 are not bases but salts
containing the bases OH− and NH2-. Also, some substances,
which many chemists considered to be acids, such
as SO3 or BCl3, are excluded from this classification due to lack
of hydrogen. Gilbert Lewis wrote in 1938, "To restrict the group
of acids to those substances that contain hydrogen interferes
as seriously with the systematic understanding of chemistry as
would the restriction of the term oxidizing agent to substances
containing oxygen."
Lewis definition:
The hydrogen requirement of Arrhenius and Brønsted–Lowry
was removed by the Lewis definition of acid–base reactions,
devised by Gilbert N. Lewis in 1923, in the same year as
Brønsted–Lowry, but it was not elaborated by him until
1938. Instead of defining acid–base reactions in terms of
protons or other bonded substances, the Lewis definition
defines a base (referred to as a Lewis base) to be a compound
that can donate an electron pair, and an acid (a Lewis acid) to
be a compound that can receive this electron pair.
In this system, an acid does not exchange atoms with a base,
but combines with it. For example, consider this classical
aqueous acid–base reaction:
HCl (aq) + NaOH (aq) → H2O (l) + NaCl (aq)
The Lewis definition does not regard this reaction as the
formation of salt and water or the transfer of H+ from HCl
to OH−. Instead, it regards the acid to be the H+ ion itself, and
the base to be the OH− ion, which has an unshared electron
pair. Therefore, the acid–base reaction here, according to the
Lewis definition, is the donation of the electron pair
from OH− to the H+ ion. This forms a covalent bond
between H+ and OH−, thus producing water (H2O).
By treating acid–base reactions in terms of electron pairs
instead of specific substances, the Lewis definition can be
applied to reactions that do not fall under other definitions of
acid–base reactions. For example, a silver cation behaves as an
acid with respect to ammonia, which behaves as a base, in the
following reaction:
Ag+ + 2 :NH3 → [H3N:Ag:NH3]+
The result of this reaction is the formation of an ammonia–
silver adduct.
In reactions between Lewis acids and bases, there is the
formation of an adduct when the highest occupied molecular
orbital (HOMO) of a molecule, such as NH3 with available lone
electron pair(s) donates lone pairs of electrons to the electrondeficient molecule's lowest unoccupied molecular orbital
(LUMO) through a co-ordinate covalent bond; in such a
reaction, the HOMO-interacting molecule acts as a base, and
the LUMO-interacting molecule acts as an acid. In highly-polar
molecules, such as boron trifluoride (BF3), the
mostelectronegative element pulls electrons towards its own
orbitals, providing a more positive charge on the lesselectronegative element and a difference in its electronic
structure due to the axial or equatorial orbiting positions of its
electrons, causing repulsive effects from lone pair – bonding
pair (Lp–Bp) interactions between bonded atoms in excess of
those already provided by bonding pair – bonding pair (Bp–Bp)
interactions. Adducts involving metal ions are referred to as coordination compounds.
Other acid–base theories:
Usanovich definition:
Mikhail Usanovich developed a general theory that does not
restrict acidity to hydrogen-containing compounds, but his
approach, published in 1938, was even more general than
Lewis theory. Usanovich's theory can be summarized as
defining an acid as anything that accepts negative species or
donates positive ones, and a base as the reverse. This pushed
the concept of acid–base reactions to its logical limits, and even
redefined the concept of redox (oxidation-reduction) as a
special case of acid-base reactions, and so did not become
widespread, despite being easier to understand than Lewis
theory, which required detailed familiarity with atomic
structure. Some examples of Usanovich acid-base reactions
include:
Na2O (base) + SO3 (acid) → 2 Na+ + SO42- (species exchanged:
anion O2−)
3 (NH4)2S (base) + Sb2S3 (acid) → 6 NH4++ 2 SbS42- (species
exchanged: anion S2−)
Na (base) + Cl (acid) → Na+ + Cl− (species exchanged: electron)
Lux–Flood definition:
This acid–base theory was a revival of oxygen theory of acids
and bases, proposed by German chemist Hermann Lux in 1939,
further improved by Håkon Flood circa 1947 and is still used in
modern geochemistry and electrochemistry of molten salts.
This definition describes an acid as an oxide ion (O2−) acceptor
and a base as an oxide ion donor. For example:
MgO (base) + CO2 (acid) → MgCO3
CaO (base) + SiO2 (acid) → CaSiO3
NO3- (base) + S2O72- (acid) → NO2+ + 2 SO42Pearson definition:
In 1963, Ralph Pearson proposed an advanced qualitative
concept known as Hard Soft Acid Base principle, later made
quantitative with help of Robert Parr in 1984. 'Hard' applies to
species that are small, have high charge states, and are weakly
polarizable. 'Soft' applies to species that are large, have low
charge states and are strongly polarizable. Acids and bases
interact, and the most stable interactions are hard–hard and
soft–soft. This theory has found use in organic and inorganic
chemistry.
Acid–alkali reaction:
An acid–alkali reaction is a special case of an acid–base
reaction, where the base used is also an alkali. When an acid
reacts with an alkali it forms a metal salt and water. Acid–alkali
reactions are also a type of neutralization reaction.
In general, acid–alkali reactions can be simplified to
OH−(aq) + H+(aq) → H2O
by omitting spectator ions.
Acids are in general pure substances that contain hydrogen
ions (H+) or cause them to be produced in solutions.
Hydrochloric acid (HCl) and sulfuric acid (H2SO4) are common
examples. In water, these break apart into ions:
HCl → H+(aq) + Cl−(aq)
H2SO4 → H+(aq) + HSO4- (aq)
An alkali is a base that contains a metal from column 1 or 2 of
the periodic table (the alkali metals or the alkaline earth
metals). Alkalis may be defined as soluble bases, which means
they must be able to dissolve in water. In general, bases are
defined as substances that contain hydroxide ion (OH−) or
produce it in solution. Therefore, one may also speak of
hydroxide bases that dissolve in water, and thus these would
also be alkalis. Some examples, then, of alkalis would
be sodium hydroxide (NaOH), potassium
hydroxide (KOH), magnesium hydroxide(Mg(OH)2), and calcium
hydroxide (Ca(OH)2). Note that only hydroxides with an alkali
metal — column 1 — are very soluble in water; hydroxides with
an alkaline earth metal — column 2 — are not as soluble. Some
sources will even say the alkaline earth metal hydroxides are
insoluble.
To produce hydroxide ions in water, the alkali breaks apart into
ions as below:
NaOH → Na+(aq) + OH−(aq)
However, alkalis may also have a broader definition that
includes carbonates (CO32-) bonded to a column 1 metal, an
ammonium ion (NH4+), or an amine (NHx radical) as the positive
ion. Examples of alkalis would then also include Li2CO3, Na2CO3,
and (NH4)2CO3.
There are many uses of neutralization reactions that are acidalkali reactions. A very common use is antacid tablets. These
are designed to neutralize excess stomach acid (HCl) that may
be causing discomfort in the stomach or lower esophagus. Also
in the digestive tract, neutralization reactions are used when
food is moved from the stomach to the intestines. In order for
the nutrients to be absorbed through the intestinal wall, an
alkaline environment is needed, so the pancreas produce an
antacid bicarbonate to cause this transformation to occur.
Another common use, though perhaps not as widely known, is
in fertilizers and control of soil pH. Slaked lime (calcium
hydroxide) or limestone (calcium carbonate) may be worked
into soil that is too acidic for plant growth. Fertilizers that
improve plant growth are made by neutralizing sulfuric acid
(H2SO4) or nitric acid (HNO3) with ammonia gas (NH3),
making ammonium sulfate or ammonium nitrate. These are
salts utilized in the fertilizer. Industrially, a by-product of the
burning of coal, sulfur dioxide gas may combine with water
vapor in the air to eventually produce sulfuric acid, which falls
as acid rain. To prevent the sulfur dioxide from being released,
a device known as a scrubber gleans the gas from smoke stacks.
This device first blows calcium carbonate into the combustion
chamber where it decomposes into calcium oxide (lime) and
carbon dioxide. This lime then reacts with the sulfur dioxide
produced forming calcium sulfite. A suspension of lime is then
injected into the mixture to produce a slurry, which removes
the calcium sulfite and any remaining unreacted sulfur dioxide.
By : Afrah Eid AL.azmi